Electrochemistry and Batteries: Complete Beginner's Guide

Electrochemistry studies the two-way street between chemical reactions and electricity. In a galvanic cell, a spontaneous reaction produces current, which is how batteries work. In an electrolytic cell, current drives a non-spontaneous reaction, which is how electroplating works.

It matters because the world runs on these cells. Every phone and electric vehicle carries a battery, corrosion eats huge amounts of steel yearly through electrochemical attack, and green hydrogen may be produced by electrolyzing water. The Nernst equation and Faraday's laws quantify it all.

Key Concepts

• Galvanic (voltaic) cell: spontaneous redox reaction split into half-cells produces current. Example: the Daniell cell, zinc and copper.
• Electrolytic cell: an external power supply forces a non-spontaneous reaction, as in copper electroplating.
• Anode and cathode: oxidation always happens at the anode, reduction at the cathode. In a galvanic cell the anode is negative; in an electrolytic cell it is positive.
• Salt bridge: completes the circuit with ions; electrons travel through the external wire, never the bridge.
• Standard reduction potential (E naught): each half-reaction's voltage versus the hydrogen electrode, in volts. Tables rank oxidizing strength.
• Nernst equation: corrects the cell voltage for non-standard concentrations.
• Faraday's laws: the mass deposited or dissolved is proportional to the total charge passed.

Formulas

• Cell EMF: E cell = E cathode minus E anode, using reduction potentials. Positive means spontaneous.
• Nernst equation: E = E naught minus (R T / n F) times ln Q. At 25 C this becomes E = E naught minus (0.0592 / n) times log10 Q, with n electrons transferred.
• Faraday's first law: mass m = (Q times M) divided by (n times F), where Q is charge in coulombs, M molar mass in g/mol, and F = 96485 C/mol.
• Charge: Q = I times t, current in amperes times time in seconds.
• Gibbs link: delta G = minus n F E; a positive cell voltage means negative free energy, hence spontaneous.

Worked Example

A Daniell cell has E naught = 1.10 V (n = 2). (a) Find E when [Zn2+] = 0.10 M and [Cu2+] = 1.0 M. (b) How much copper plates out if 2 A flows for 30 minutes?

1. Reaction quotient Q = [Zn2+]/[Cu2+] = 0.10/1.0 = 0.10.
2. E = 1.10 minus (0.0592/2) times log10(0.10) = 1.10 minus 0.0296 times (minus 1) = 1.1296 V, about 1.13 V.
3. Lower product concentration raises the voltage, exactly as Le Chatelier predicts.
4. Charge Q = 2 A times 1800 s = 3600 C. Moles of electrons = 3600/96485 = 0.0373 mol.
5. Copper needs 2 electrons per atom: 0.0187 mol Cu times 63.55 g/mol = 1.19 g of copper plated.

Common Misconceptions

• Electrons flow through the salt bridge. They do not: electrons travel through the external wire, while ions migrate through the bridge to balance charge.
• The anode is always negative. Anode means oxidation site. It is negative in a galvanic cell but positive in an electrolytic cell, where the supply drags electrons out of it.
• A bigger battery voltage means it stores more energy. Voltage sets the push; capacity in ampere-hours sets the stored charge. Energy needs both.

Real-World Applications

• Batteries: lead-acid car batteries, lithium-ion phones and EVs, all galvanic cells refined by electrochemistry.
• Electroplating: chrome bumpers, gold connectors, and corrosion-resistant coatings deposited by Faraday's laws.
• Corrosion protection: sacrificial zinc anodes and cathodic protection guard ships and pipelines.
• Metal refining: aluminium and copper are purified by electrolysis on an industrial scale.

Split a redox reaction into half-cells and you get a battery; push current backward and you get plating and refining. The Nernst equation tunes voltage to real concentrations and Faraday's laws convert charge into grams.